Chapter 1 Review61 Total Terms
Chapter 1 Key Terms & Vocabulary
Comprehensive index of all scientific terms, definitions, and examples across Chapter 1: Essential Ideas.
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From Section 1.1: Chemistry in Context
chemistryGeneral Chemistry
The study of the composition, properties, and interactions of matter.
Example: Synthesizing polymers or refining petroleum.
From Section 1.1: Chemistry in Context
hypothesisScientific Method
A tentative explanation of observations that acts as a guide for gathering and checking information.
Example: Proposing that a plant extract lowers blood pressure.
From Section 1.1: Chemistry in Context
lawsScientific Method
Statements that summarize a vast number of experimental observations, describing or predicting facets of the natural world.
Example: Law of Conservation of Mass.
From Section 1.1: Chemistry in Context
theoriesScientific Method
Well-substantiated, comprehensive, testable explanations of particular aspects of nature.
Example: Dalton
From Section 1.1: Chemistry in Context
scientific methodScientific Method
The path of discovery leading from question and observation to law or hypothesis to theory, backed by experimental verification.
Example: Testing a proposed medical treatment through controlled clinical trials.
From Section 1.1: Chemistry in Context
macroscopic domainChemistry Domains
The realm of everyday things that are large enough to be sensed directly by human sight or touch.
Example: Observing ice melting or water boiling.
From Section 1.1: Chemistry in Context
microscopic domainChemistry Domains
The realm of atoms, molecules, and subatomic particles that are typically too small to be seen directly.
Example: Visualizing individual water molecules or chemical bonds.
From Section 1.1: Chemistry in Context
symbolic domainChemistry Domains
The specialized language used to represent components of the macroscopic and microscopic domains.
Example: Chemical formulas like H₂O and reaction equations.
From Section 1.2: Phases and Classification of Matter
MatterGeneral
Anything that occupies space and has mass.
Example: Water, rocks, and air.
From Section 1.2: Phases and Classification of Matter
solidStates of Matter
State of matter characterized by rigidity and a definite shape and volume.
Example: An ice cube or diamond.
From Section 1.2: Phases and Classification of Matter
liquidStates of Matter
State of matter that flows, takes the shape of its container, and has a fixed volume.
Example: Water or liquid mercury.
From Section 1.2: Phases and Classification of Matter
gasStates of Matter
State of matter that takes both the shape and volume of its container.
Example: Oxygen or helium.
From Section 1.2: Phases and Classification of Matter
plasmaStates of Matter
Gaseous state of matter containing an appreciable number of electrically charged particles.
Example: Stars and lightning.
From Section 1.2: Phases and Classification of Matter
massProperties
A measure of the amount of matter in an object.
Example: Measured in grams or kilograms.
From Section 1.2: Phases and Classification of Matter
WeightProperties
The force that gravity exerts on an object, directly proportional to its mass.
Example: An object weighs less on the Moon.
From Section 1.2: Phases and Classification of Matter
law of conservation of matterFundamental Laws
States that there is no detectable change in total mass during physical or chemical changes.
Example: Total mass of reactants equals total mass of products.
From Section 1.2: Phases and Classification of Matter
pure substanceClassification
Matter that has a constant composition and identical properties throughout.
Example: Pure water or table sugar.
From Section 1.2: Phases and Classification of Matter
elementsClassification
Pure substances that cannot be broken down into simpler substances by chemical changes.
Example: Gold, silver, and oxygen.
From Section 1.2: Phases and Classification of Matter
compoundsClassification
Pure substances comprised of two or more elements chemically combined in fixed ratios.
Example: Water (H₂O) or sodium chloride.
From Section 1.2: Phases and Classification of Matter
mixtureClassification
Matter composed of two or more types of matter present in varying amounts and separable physically.
Example: Trail mix or salt water.
From Section 1.2: Phases and Classification of Matter
heterogeneous mixtureClassification
A mixture with a composition that varies from point to point.
Example: Italian dressing or granite.
From Section 1.2: Phases and Classification of Matter
homogeneous mixtureClassification
A mixture that exhibits a uniform composition and appears visually the same throughout.
Example: Sports drink or air.
From Section 1.2: Phases and Classification of Matter
solutionClassification
Another name for a homogeneous mixture, uniform at the molecular level.
Example: Salt dissolved in water.
From Section 1.2: Phases and Classification of Matter
atomAtomic Theory
The smallest particle of an element that has the properties of that element and enters into chemical combinations.
Example: A single gold atom.
From Section 1.2: Phases and Classification of Matter
DaltonHistory
John Dalton, the British scientist who established modern quantitative atomic theory.
Example: Proposed atomic theory in the early 19th century.
From Section 1.2: Phases and Classification of Matter
moleculeAtomic Theory
Consists of two or more atoms joined by strong chemical bonds acting as a single unit.
Example: Water molecule (H₂O) or oxygen gas (O₂).
From Section 1.3: Physical and Chemical Properties
Physical PropertyProperties of Matter
A characteristic of matter that is not associated with any change in its chemical composition.
Example: Density, color, hardness, and melting point.
From Section 1.3: Physical and Chemical Properties
Physical ChangeChanges in Matter
A change in the state or appearance of matter without any accompanying change in its chemical identity.
Example: Melting ice, dissolving sugar, or condensing steam.
From Section 1.3: Physical and Chemical Properties
Chemical PropertyProperties of Matter
The ability or inability of a substance to change into one or more other types of matter.
Example: Flammability, toxicity, acidity, and iron
From Section 1.3: Physical and Chemical Properties
Chemical ChangeChanges in Matter
A change that produces one or more types of matter that differ entirely from the matter present before the change.
Example: Combustion (burning), cooking food, and metal rusting.
From Section 1.3: Physical and Chemical Properties
Extensive PropertyProperty Classification
A property of matter whose value depends directly on the amount of substance present.
Example: Mass and volume.
From Section 1.3: Physical and Chemical Properties
Intensive PropertyProperty Classification
A property of matter whose value is independent of the amount of substance present.
Example: Temperature and density.
From Section 1.4: Measurements
UnitsMeasurement
Standards of comparison used to give meaning to numerical measurements.
Example: Liters, grams, meters
From Section 1.4: Measurements
International System of UnitsStandards
Globally agreed-upon system of seven base physical units, abbreviated as SI.
Example: The SI system used in scientific research worldwide.
From Section 1.4: Measurements
SI unitsStandards
Standardized units established by international agreement (Le Système International d’Unités).
Example: Kilograms, meters, seconds
From Section 1.4: Measurements
lengthProperties
Measure of one dimension of an object or distance between points.
Example: Measuring the height of a beaker in meters.
From Section 1.4: Measurements
meter (m)SI Units
The standard SI base unit of length.
Example: 1 m = 39.37 inches
From Section 1.4: Measurements
kilogram (kg)SI Units
The standard SI base unit of mass.
Example: 1 kg is approximately 2.2 pounds.
From Section 1.4: Measurements
IUPACOrganizations
International Union of Pure and Applied Chemistry; sets international chemical standards.
Example: Redefined the kilogram standard in 2019.
From Section 1.4: Measurements
kelvin (K)SI Units
The standard SI base unit of temperature, with zero representing absolute zero.
Example: Water freezes at 273.15 K.
From Section 1.4: Measurements
Celsius (°C)Temperature
Commonly used temperature scale where water freezes at 0 °C and boils at 100 °C.
Example: Normal human body temperature is 37 °C.
From Section 1.4: Measurements
second (s)SI Units
The standard SI base unit of time.
Example: Measuring the duration of a chemical reaction.
From Section 1.4: Measurements
VolumeProperties
The measure of the amount of space occupied by an object.
Example: A 2-liter soft drink bottle.
From Section 1.4: Measurements
cubic meter (m3)Derived Units
The standard SI derived unit of volume, representing a cube with 1 m edges.
Example: m³
From Section 1.4: Measurements
liter (L)Derived Units
A common unit of volume equal to one cubic decimeter (dm³).
Example: A 1-L graduated cylinder.
From Section 1.4: Measurements
cubic centimeter (cm3)Derived Units
Volume of a cube with 1 cm edge length, equal to one milliliter.
Example: cm³
From Section 1.4: Measurements
milliliter (mL)Derived Units
Volume unit equal to one-thousandth of a liter (10⁻³ L).
Example: 50 mL of hydrochloric acid.
From Section 1.4: Measurements
densityProperties
The ratio of an object
Example: Gold has a density of 19.3 g/cm³.
From Section 1.4: Measurements
ccDerived Units
Cubic centimeter; abbreviation frequently used in medical settings.
Example: A dose of 5 cc of medicine.
From Section 1.5: Measurement Uncertainty, Accuracy, and Precision
exact numberMeasurements
A value known with complete certainty, obtained through counting or definition.
Example: 1 foot = 12 inches
From Section 1.5: Measurement Uncertainty, Accuracy, and Precision
uncertaintyMeasurements
A quantitative estimate of how much a measured value deviates from the actual result.
Example: 6.72 g ± 0.01 g
From Section 1.5: Measurement Uncertainty, Accuracy, and Precision
significant figuresData Analysis
All digits in a measurement known with certainty plus one final estimated digit.
Example: 21.6 mL has 3 significant figures
From Section 1.5: Measurement Uncertainty, Accuracy, and Precision
significant digitsData Analysis
Alternative terminology for significant figures reflecting measurement certainty.
Example: 4.50 has 3 significant digits
From Section 1.5: Measurement Uncertainty, Accuracy, and Precision
roundingMath Operations
The procedure of reducing digits in a number while keeping its value close, using standard tie-breaking rules.
Example: 0.02867 rounds to 0.0287
From Section 1.5: Measurement Uncertainty, Accuracy, and Precision
precisionExperimental Quality
The degree of agreement among repeated measurements of the same quantity.
Example: Getting 10.1 mL, 10.2 mL, and 10.1 mL on repeated trials
From Section 1.5: Measurement Uncertainty, Accuracy, and Precision
accuracyExperimental Quality
The degree of closeness between a measured value and the true or accepted reference value.
Example: A measured density of 7.8 g/cm³ compared to a true value of 7.9 g/cm³
From Section 1.6: Mathematical Treatment of Measurement Results
dimensional analysisCalculation Methods
A mathematical approach where units of quantities are subjected to the same mathematical operations as numbers to solve problems.
Example: Converting speed from m/s to km/h
From Section 1.6: Mathematical Treatment of Measurement Results
factor-label methodCalculation Methods
Alternative name for dimensional analysis, emphasizing the tracking and cancellation of unit labels.
Example: Multiplying by (1 m / 100 cm)
From Section 1.6: Mathematical Treatment of Measurement Results
unit conversion factorMeasurement
A ratio of two equivalent quantities expressed with different measurement units, equal in value to 1.
Example: 1 kg / 1000 g
From Section 1.6: Mathematical Treatment of Measurement Results
temperatureThermal Physics
A quantitative measure of the hotness or coldness of a substance.
Example: Boiling point of water is 100 °C
From Section 1.6: Mathematical Treatment of Measurement Results
FahrenheitThermal Physics
A temperature scale where water
Example: Normal body temperature is 98.6 °F